CATALYSIS  IN  THE  DECOMPOSITION 
OF  POTASSIUM  CHLORATE 


BY 


MYRON  ALONZO  SNELL 


THESIS 


FOR  THE 


DEGREE  OE  BACHELOR  OF  SCIENCE 


CHEMISTRY 


COLLEGE  OF  LIBERAL  ARTS  AND  SCIENCES 


UNIVERSITY  OF  ILLINOIS 


1922 


Digitized  by  the  Internet  Archive 
in  2015 


https://archive.org/details/catalysisindecomOOsnel 


) 922 
Sh  2 


UNIVERSITY  OF  ILLINOIS 


l-lay__3_5 


--IQ2A 


THIS  IS  TO  CERTIFY  THAT  THE  THESIS  PREPARED  UNDER  MY  SUPERVISION  BY 

_ _Al_onzo_  _Snell_  _ 

entitled C at  I s_  _ In  _ jbhe_  JDe  oo  si  t_i  on  _ o f _ Pot  as  siuin 

__Chlorate_. 


IS  APPROVED  BY  ME  AS  FULFILLING  THIS  PART  OF  THE  REQUIREMENTS  FOR  THE 
DEGREE  OF  P£_3_Cc£_*v.C,e* 

d»v__C_  RenUwSTtxa 


_Cj_ 

Instructor  in  Charge 


Approved 


HEAD  OF  DEPARTMENT  OF 


TABLE  OF  CONTENTS. 


Aokno  wl e dgrae n t 

I.  Introduction 

II.  Historical 

III.  theoretical 

The  Course  of  the  Reaction 

IV.  Experimental 

a.  Materials 

b.  Apparatus 

V.  Experimental  Results 

a.  Qualitative  Tests 

b.  Nature  of  the  Catalyst 

c.  Relative  Activity  of  Catalysts 

d.  Effect  of  Varying  Percent  of  Catalyst 

e.  Promoter  Action 
VI.  Summary 

VII.  Bibliography 


Page 

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3 

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6 

7 

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9 

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11 

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13 

33 


- 1 ~ 


I 

INTRODUCTION 

The  decomposition  of  potassium  chlorate  is  facilitated  by 
certain  substances  which  act  as  catalysts  in  the  reaction.  The 
catalytic  effect  includes  a considerable  lowering  of  the  temper- 
ature at  which  the  decomposition  begins,  and  a tremendous  increase 
in  the  rate  at  which  the  oxygen  is  given  off*  When  potassium 
chlorate  is  heated  alone,  oxygen  begins  to  come  off  at  about 
340o-350®,  and  manganese  dioxide,  heated  alone,  gives  off  oxygen, 
beginning  at  about  400®  (Alex  Smith),  but  with  a mixture  of  the 
two,  the  evolution  commences  at  about  225®,  more  than  100®  lower. 
The  rate  of  evolution  of  the  gas  increases  very  slowly  as  the 
temperature  is  raised  until,  at  305®,  it  suddenly  becomes  rapid, 
and  at  about  335®  as  the  chlorate  begins  to  melt,  it  is  toe  great 
to  be  measured. 


< 


- 


. 


- 2 - 
II 

HISTORICAL 

It  is  not  known  who  was  the  first  chemist  to  observe  that 
certain  catalytic  agents  accelerate  the  decomposition  of  potass 
ium  chlorate  to  give  oxygen,  but  several  men  have  added  some- 
thing to  the  knowledge  of  the  subject  in  the  last  seventy- 
five  years.  The  first  contribution  was  that  of  Wachter,  in 
1843.  Other  prominent  contributors  were  Baudrimont  and 
Jungfleisch  in  1871;  Schulze  in  1860;  Hodgkinson,  Lowndes, 
and  Veley  in  1888-89;  Fowler  and  Grant  in  1890;  McLeod  in 
1889,  189  4 and  1896;  Brunck  in  1893;  and  Sodeau  in  1900; 

1901,  and  1902.  (The  mo3t  important  results  of  these  investi- 
gators is  outlined  below.)  Nothing  much  of  importance 
has  been  done  on  it  since  the  work  of  Sodeau. 


. 


. 


- 3 - 
III 

THEORETICAL 

The  Course  of  the  Reaction. 

A great  many  different  reactions  have  been  suggested  for 
the  decomposition  and  there  still  exists  a great  deal  of 
disagreement  on  this  question. 

McLeod  in  1896  considered  the  decomposition  to  be  a 
chemical  reaction,  which  involved  the  catalyst  (approximate 
reactions) : 

2 KC103  + 2 MnOs  * 3 KMn04  4 Cl8  + Og 
2 KMn04  * K8Mn04  + MnOa  + 08 

KgMn04  + Cl*  * 2 KC1  + MnOa  + Og 

He  found,  also,  that  the  particles  of  MnOg  are  broken 
up  during  the  decomposition  by  examining  them  with  a micro- 
scope. Baudrimont  and  Veley  claim  than  an  evolution  of 
oxygen  can  be  produced  without  fusing  the  chlorate  and  with 
no  catalyst  present.  And  since  many  inert  substances,  such 
as  Si02  do  not  have  a catalytic  effect,  Soaeau  stated  that 
what  actually  took  place  was  probably  a chemical  reaction 
involving  the  catalyst,  rather  than  a physical  efxect  due 
to  contact. 

Sodeau  also  reasoned  that  the  chlorate  probably  under- 
goes self-oxidat ion  when  heated  alone,  and  the  perchlorate 


* 

. 

» 


is  formed: 


- 4 - 


4 KC103  + 61,300  cal.  * KC1  4 3KC104 
His  idea  is  confirmed  by  the  following  facts: 

4 KClOa  - 39,000  cal.,  4 KC1  + fo 


4 KCXO3  cal.>  3 KCIO*  4 KC1 


Then  by  Hess's  law: 

3 KC10*  4 KC1  l~i00»500  Cal->4  KOI  46  0* 

Since  any  react  ion ^will  always  take  place  in  the  direction 
in  which  heat  is  absorbed,  the  above  reaction  will  go  through 
the  formation  and  decomposition  of  the  perchlorate  at  the 
expense  of  the  other  one.  Consequently  unless  a catalyst  is 
present,  the  perchlorate  is  always  formed  as  an  intermediate 
compound.  This  reaction  occurs  only  at  a temperature  of 
340®  or  higher. At  this  temperature,  only  a small  amount  of  the 
total  oxygen  present  in  the  ohlorate  is  given  off,  as  shown 
by  Remsen's  equation,  and  confirmed  by  experimental  results: 

8 KC10a  = 5 KCIO4  + 3 KC1  + 3 0S 
But  when  a catalyst  is  present,  no  perchlorate  is  formed, 
even  temporarily  from  the  chlorate,  according  to  Sodeau,  since 


1> 


t 


- 5 - 

Ecoles  found  none  in  the  residue  on  analysis,  and  McLeod 
showed  that  if  it  was  once  formed,  it  would  not  be  decomposed 
under  the  conditions  of  the  experiment,  so  that  the  reaction 
is  probably 

3 KC103  = 2 KC1  + 3 0g  (Melior) 
or  one  similar  to  it. 


- 6 - 
IV 

EXPERIMENTAL 

a*  Materials. 

C.  P.  KC103  was  used  in  all  of  the  determinations.  The 
crystals  were  very  fine  and  white.  The  commercial  MnOa  used 
was  found  on  analysis  to  contain  91.07$  MnOs,  and  6.83$ 

Fe203.  The  pure  MnOs  was  prepared  after  the  method  of 
Georgen.  C.  P.  Manganese  carbonate  was  added  in  excess  to 
nitric  acid.  The  crystals  of  Mnjho^were  washed  and  changed 
to  MnOa  by  heating  at  300®  . The  other  catalysts  employed 
were  the  impure  compounds: 

A charge  was  prepared  as  follows:  The  KC103  and  catalyst 

in  weighed  proportions,  were  mixed  thoroughly,  ground  in  a 
porcelain  mortar, moistened  with  distilled  water  and  dried.  In 
this  way,  the  catalyst  was  deposited  on  the  surface,  and  in 
the  pores  of  the  KC103  in  a finely  divided  form,  and  in 
intimate  contact  with  it.  A charge  consisted  of  6 g.  of  KC103, 
plus  an  amount  of  the  catalyst  calculated  as  a certain  percent-* 
age  of  the  total  weight.  For  example,  in  the  table,  KC103 
plus  25$  of  MnOg,  refers  to  6 g.  of  the  chlorate  mixed  with 
2 g.  of  the  dioxide. 


- 7 - 

to*  Apparatus. 

The  apparatus  used  to  study  the  decomposition  was  essen- 
tially as  shown  in  Fig.  1.  The  reaction  tutoe  was  a Pyrex 
glass  test-tube  3.5  toy  30.5  cm.  The  charge  occupied  a length 
of  from  3 to  3 cm.  in  the  bottom  of  the  tutoe.  The  flask  was 
also  of  Pyrex  glass  5 toy  35  cm.  The  gas  was  collected  toy 
downward  displacement  of  water  in  a 100  cc . eudiometer.  The 
furnace  was  connected  to  110  A.U.  terminals,  the  current 
toeing  cut  down  to  a suitable  amount  toy  means  of  a cartoon 
resistance.  The  temperature  was  measured  with  a 360°  thermom- 
eter, with  the  bulb  placed  in  the  reacting  mixture.  The 
temperature  of  the  charge  was  kept  constant  toy  means  of  a con- 
stant boiling  liquid  in  the  bottom  of  the  flask.  Benzoic 
acid  was  used  for  a temperature  of  335©,  acetanilid  for  300<> , 
and  anthracene  for  338° . The  variation  in  temperature  was 


. 


- 8 - 
V 

EXPERIMENTAL  RESULTS* 
a*  Qualitative  tests. 

The  temperatures  chosen  were  235©,  300©,  and  328©  • 

Below  235©,  the  decomposition  was  too  slight  to  he  appreciable, 
and  above  328©,  too  violent  to  be  controlled  or  measured.  The 
boiling  points  and  suitability  of  the  liquids  also  had  to  be 
considered. 

As  a basis  for  comparison,  a charge  was  run  without  any 
catalyst  present.  This  was  followed  by  experiments  on  silica, 
sodium  chloride,  sodium  sulfate,  sodium  carbonate,  alumina, 
cobalt  sulfate,  zinc  oxide,  molybdenum  oxide,  stannic  oxide, 
barium  dioxide,  bismuth  oxide,  and  mercuric  oxide,  none  of 
which  exerted  a catalytic  influence  on  the  reaction. 

Cr203,  Coa03,  and  Ni203  each  caused  the  chlorate  to  give 
off  both  oxygen  and  chlorine  at  a tremendous  rate,  even  below 
235©,  due  to  the  fact  that  they  reacted  with  the  chlorate, 
liberating  part  of  the  oxygen  and  all  of  the  chlorine.  The 
probable  reaction  for  the  Cr203  is  given  below.  K2Cr207  was 
tested  for,  and  found  in  the  residue. 

Cr203  + 2 KCiO^  — K2Cr203  4-  Cl2  + 02 
CuO,  MgO,  Pr407,  Nd203,  Mn02,  and  Fe203  ail  facil- 
itated the  reaction,  but  the  last  two,  only,  showed  a consider- 


- e - 

able  catalytic  effect,  and  were  the  only  ones  that  were 
studied  in  detail*  The  results  are  shown  in  the  accompanying 
table*  The  data  on  MnOa  and  Fe803  is  represented  graphically 
in  figures  2, 3, 4, 5, 6,7  and  8.  The  effect  of  varying  the 
temperature,  and  the  amount  of  catalyst  is  clearly  indicated 
in  these  curves.  The  rate  of  the  evolution  of  oxygen  was 
measured  in  cubic  centimeters  per  minute  at  22®,  and  760  mm* 
pressure  • 

b*  Nature  of  the  catalyst. 

Sodeau  showed  that  any  reaction  other  than  one  including 
alternate  oxidation,  and  deoxidation  of  the  catalyst  would 
involve  setting  free  of  all  of  the  chlorine,  and  its  subsequent 
complete  reabsorption,  and  since  the  amount  of  chlorine  given 
off  is  very  small,  the  catalyst  is  the  oxygen  carrier.  An 
intermediate  compound  is  probably  formed  between  the  catalyst 
and  the  oxygen  which  is  a surface  complex  of  indefinite  and 
variable  composition.  This  substance  is  unstable  and  is  almost 
immediately  broken  up  again  into  the  catalytic  agent  and  oxygen. 
The  only  substances  which  could  act  as  catalysers,  would  be 
then,  oxides  of  metals,  capable  of  more  than  one  valence.  This 
was  actually  found  to  be  the  case,  since  the  oxides  of  copper, 
magnesium,  praseodymium,  neodymium,  manganese,  and  iron  were 


. 

' T < 


- 10  - 


the  only  catalysers,  and  all  of  these  are  oxides  of  metals 
with  at  least  two  valences.  There  seems  to  be  some  doubt  as 
to  the  existence  of  another  oxide  of  magnesium  besides  MgO, 
but  R030Q3  and  Schorlemmer  state  that  Mg02  exists,  but  is 
unstable.  Magnesium  would  not,  then,  be  an  exception.  How- 
ever, there  are  metals  which,  it  would  seem,  ought  to  catalyse 
the  reaction,  such  as  barium  and  calcium,  both  of  which  form 
two  oxides.  For  some  unknown  reason  such  metals  are  exceptions 
to  the  rule. 

c.  Relative  Activity  of  Catalysts. 

The  table  shows  that  Mn02  and  Fe203  are  the  only  efficient 
catalysts  for  the  reaction,  and  that  Mn02  is  almost  twice  as 
active  as  Fe203  in  increasing  the  rate  of  gas  evolution. 
Praseodymium  oxide  i3  next  most  active,  then  the  oxides  of 
copper,  magnesium,  and  neodymium  in  the  order  mentioned. 
Manganese  dioxide  seemed  to  have  the  greatest  effect  in  lower- 
ing the  temperature  of  decomposition.  With  Mn02,  the  oxygen 
began  to  come  off  at  about  235®,  with  Fe203,  the  evolution 
began  at  a trifle  higher  temperature,  and  so  on  for  the  others. 
So  the  temperature  at  whioh  the  decomposition  begins  is  in*--- 
veraely  proportional  to  the  activity  of  the  catalyst  in 
question,  in  increasing  the  rate  of  gas  evolution. 


11 


d*  Effect  of  Varying  Percent  of  Catalyst* 

The  effect  of  varying  the  amount  of  catalyst  is  shown 
clearly  in  Figures  4,  5,  6 and  7 for  the  three  temperatures 
235©,  300©,  328©*  The  optimum  concentration  for  both  MnOa 
and  Fe203  was  about  33$ • Experiment  showed  also  that  the 
optimum  concentration  did  not  vary  with  the  temperature,  and 
that  the  amount  of  Oatalyst  present  had  no  effect  in  lowering 
the  temperature  of  decomposition* 

e*  Promoter  Action. 

Impure  MnOs  was  found  to  be  more  active  than  the  pure 
MnOg*  Analysis  showed  this  to  be  due  to  the  presence  of  some 
Fe203  as  an  impurity*  So  a charge  was  made  up  of  22*8$  pure 
MnOs  and  2.2$  pure  Fe203,  the  proportions  existing  in  the 
impure  Mn02*  As  shown  in  the  table,  this  mixture  gave  essen- 
tially  the  same  results  as  did  the  impure  Mn02.  This  is  evi- 
dence for  the  belief  that  the  Fe203  exerts  a promoter  action 
on  the  MnOa  in  catalysing  the  reaction  and  vice  versa.  This 
action  was  studied  with  various  concentrations  of  both  catalysts 
at  the  three  different  temperatures.  The  results  are  shown 
graphically  in  Figure  8*  The  optimum  concentration  was  found 
to  be  50$  of  each  catalyst*  Thi3  mixture  was  much  more  active 
than  100$  of  either  one. 


- 12  - 
VI 

SUMMARY 

1.  A number  of  s^ustances  have  been  tested  for  a catalytic 
effect  upon  the  decomposition  of  pota33ium  chlorate.  Only 
certain  metallic  oxides  are  active  in  this  respect. 

2.  The  two  best  accelerating  agents,  MfiOg  and  Fe203  have  been 
studied  at  several  temperatures,  and  in  varying  concentra- 
tions. The  optimum  concentration  of  the  catalyst  has  been 
found  to  be  approximately  33 $ by  wei^it. 

3.  Promoter  action  is  indicated  by  the  fact  that  the  activity 

of  Mn02  is  improved  by  an  admixture  of  Fe2Q3,  and  vice  versa* 
The  superior  activity  of  pyrolusite  to  pure  MnOs  is  thus 
explained  by  the  presence  of  Fe203  as  an  impurity  in  the 
ore  • 

4.  A sudden,  abnormal  increase  in  the  rate  of  the  evolution 
of  oxygen  at  about  305©  (as  illustrated  in  the  curves) 
indicates  that  this  is  a critical  point  in  the  de comp os it ion » 

5.  It  appears  from  available  data  that  the  effect  of  heating 
KC103  alone  may  be  expressed  by  two  concurrent  reactions: 

(1)  4 KC103  = 3 KC104  * KC1 

(2)  2 KC103  = KC1  + 3 08 


of  which  the  first  is  of  much  greater  magnitude,  especially 
if  the  temperature  is  low.  The  XC104  subsequently  decom- 
poses at  a higher  temperature  into  KC1  and  oxygen. 


- 13  - 


6*  In  the  presence  of  a catalytic  agen$,  the  de  compos  it  ion 
of  KC103  proceeds  by  the  second  equation,  at  the  expense 
of  the  first,  or  more  probably,  by  an  entirely  different 
route,  in  which  an  intermediate  compound  is  forma* 
between  oxygen  and  the  catalyst*  This  compound  is  un- 
stable and  is  probably  of  indefinite  and  variable  compo- 
sition - of  the  nature  of  a surface  complex.  This  con- 
clusion would  seem  to  be  supported  by  the  fact  that  the 
only  substances  which  are  active  accelerating  agent 3 are 
the  oxides  of  metals  capable  of  more  than  one  valence* 


- 14- 


Material 
$ Catalyst 

335° 

CC.  02  per  minute 
300° 

328° 

KClOs  (6  g. ) 

— 

0.35 

1.0 

7.7  MnO 2 

1.9 

3.  55 

33.5 

14.3  MnO  g 

3.55 

5.0 

59.0 

35.0  Mn02 

5. 95 

8.0 

94,5 

39. 4 MnO 2 

7.8 

8.6 

113. Q 

33.3  MnO 2 

8.3 

13.4 

119.5 

40.0  MnO g 

8.1 

13.0 

115.5 

50.0  MnO 2 

7.5 

11.5 

104.3 

1 4.  3 Fe  gO 3 

2.1 

3.3 

34.0 

3a. 0 Fe gO 3 

3.6 

5.3 

51 . 5 

35,3  Fe 2O3 

4.85 

6.65 

62.0 

4 a ,0  £ e g 9 3 

4.3 

6.4 

59. 0 

a 0 . 0 Fe  20  3 

3,5 

5.85 

51.0 

18.75  MnO s ) 
6.35  Fe303) 

7.0 

11.3 

99.1 

18.75  Fe 2O3) 
6.35  MnO 2 ) 

6.0 

9.5 

- - 89.0 

...  V 

13.5  MnO 2 ) 

13.5  Fe  gO  3 ) 

7.9 

13.35 

ii&.s 

35.0  CuO 

0.95 

1.95 

34.0 

35.0  Pr407 

3.35 

3.75 

38.0 

35.0  Nd303 

0045 

0.65 

5.05 

35.0  MgO 

. 45 

.8 

7.35 

35.0  MnOgCC.P.  ) 

1.8 

3,  6 

37.5 

33.8  Mn02(C.P.  ) 

5.7 

7.8 

86.  0 

• 3 F©  gQ  g ^ 


t 


1 ft 


15 


Te.  MPE  RATURE. 


Fioune  2j 


16- 


-17- 


o »o  Zo  zo  40 

Caocentratioo  tn  percent 


FtQ  U RE  4 


CC.  GAS  PER  MIN  U X E. 


-IS- 


CC.' . P£  K H I M U T t. 


~|5~ 


C e . G A C)  R r^i  I 


- 20  - 


o 


- Z I - 


22  - 


VII. 

BI3LI0GRAPHY 

Baudrimont  - J.  Pharm.  Chim.  1871  (IV)  _14,  81  and  161. 

.Jungf lelsoh  - Ibid  1871  (IV)  .14,  130. 

H.  Schulze  - J.  Pr.  Chem.  1880  (II)  21,  426. 

Hodgkinson  and  Lowndes  - Chem.  News,  1888,  J58,  309;  1889,  _59,  63 
Veley  - Phil.  Trans.  - 1888,  A . 271. 

Fowler  and  Grant  - Trans.  1890,  .57,  272. 

McLeod  - Trans.  1889,  55,  184;  1894,  .64,  202:  1896,  69,  1015. 

Brunck,-  Ber.  1893  , 26,  1790;  Zeit.  Anorg.  Chem.  1895,  10,  222. 
Wachter  - J • Pr.  Chem.  1843,  30.  325. 

Sodeau  - J.  Chenu  Soc.  1900,  77,  137  and  717;  1901,  79,  247  and 
939;  1902,  81,  1066. 

Mellor  , J.  W.  - Modern  Inorg.  Chem. 

Alex  Smith  - Inorg.  Chem. 

Remsen  - Chemistry  (1906) 

Georgen  - Compt.  rend.  1879,  J58,  797. 

Roscoe  and  Schorlemmer  - Treatise  on  Chemistry,  Vol.  2. 


